Ammonia Ammonia Ammonia The manufacture of ammonia is crucial for the world's agricultural industry for from it all fertilizers that contain nitrogen are produced. Uses of ammonia The manufacture of fertilizers is by far the most important use of ammonia. These include urea , ammonium salts (ammonium phosphates, ammonium nitrate, calcium ammonium nitrate) and solutions of ammonia. Figure 1 The uses of ammonia. An increasing amount of ammonia, although still small compared with other uses, is used as a concentrated solution in combating the discharge of nitrogen oxides from power stations. Annual production of ammonia Ammonia ranks second, to sulfuric acid, as the chemical with the largest tonnage. It is being increasingly made in countries which have low cost sources of natural gas and coal (China and Russia account for ca 40%). The largest plants produce about 3000 tonnes a day and there are plans to build plants that produce 4000-5000 tonnes a day, which would mean that the total world output could be managed with 100 such units. Current production is: World Data from: 1.

Federal State Statistics Service: Russian Federation 2011 The world population is increasing by about 1.4% a year and the increase in tonnage of ammonia made just about keeps pace. Manufacture of ammonia The manufacture of ammonia from nitrogen and hydrogen takes place in two main stages: a) the manufacture of hydrogen b) the synthesis of ammonia (the Haber Process) The manufacture of hydrogen involves several distinct processes. Figure 2 shows their sequence and the location within an ammonia plant (steps1-5). The converter used to make ammonia from the hydrogen is also shown (step 6). What occurs in each of these steps is described below the figure. Figure 2 An ammonia plant in Western Australia: 1 Desulfurisation units (a) The manufacture of hydrogen Hydrogen is produced from a variety of feedstocks, mostly from natural gas, coal or naphtha. The ways in which hydrogen is obtained from these feedstocks are dealt with separately.

Hydrogen from natural gas (methane) This involves two stages: i) the manufacture of synthesis gas (a mixture of carbon monoxide and hydrogen (steam reforming)) >ii) the removal of the carbon monoxide and production of a mixture of hydrogen and nitrogen (the shift reaction) (i) The manufacture of synthesis gas Whichever way the methane is obtained, it will contain some organic sulfur compounds and hydrogen sulfide, both of which must be removed. Otherwise, they will poison the catalyst needed in the manufacture of synthesis gas. In the desulfurisation unit, the organic sulfur compounds are often first converted into hydrogen sulfide, prior to reaction with zinc oxide. The feedstock is mixed with hydrogen and passed over a catalyst of mixed oxides of cobalt and molybdenum on an inert support (a specially treated alumina) at ca 700 K. Then the gases are passed over zinc oxide at ca 700 K and hydrogen sulfide is removed: Primary steam reforming converts methane and steam to synthesis gas, a mixture of carbon monoxide and hydrogen: High temperatures and low pressures favour the formation of the products (Le Chatelier's Principle).

In practice, the reactants are passed over a catalyst of nickel, finely divided on the surface of a calcium oxide/aluminium oxide support contained in vertical nickel alloy tubes. The tubes, up to 350 in parallel, are heated in a furnace above 1000 K and under a pressure of ca 30 atm. This is an example of a tubular reactor . Secondary steam reforming reacts oxygen from the air with some of the hydrogen present and the resulting mixture is passed over a nickel catalyst. The steam and heat produced from the combustion reforms most of the residual methane. Among the key reactions are: The emerging gas from this net exothermic stage is at ca 1200 K and is cooled in heat exchangers . The steam formed from the water used in cooling the gases is used to operate turbines and thus compressors and to preheat reactants. Some recent designs use waste heat from the secondary reformer directly to provide heat for the primary reformer. At this stage the gas contains hydrogen, nitrogen, carbon monoxide and carbon dioxide and about 0.25% methane. As air contains 1% argon, this also accumulates in the synthesis gas.

(ii) The shift reaction This process converts carbon monoxide to carbon dioxide, while generating more hydrogen. It takes place in two stages. In the first, the high temperature shift reaction, the gas is mixed with steam and passed over an iron/chromium(III) oxide catalyst at ca 700 K in a fixed bed reactor . This decreases the carbon monoxide concentration from 11%: In the second stage, the low temperature shift reaction, the mixture of gases is passed over a copper-zinc catalyst at ca 500 K. The carbon monoxide concentration is further reduced to 0.2%. The reaction is done in two stages for several reasons. The reaction is exothermic. However, at high temperature, the exit concentration of carbon monoxide is still quite high, due to equilibrium control. The copper catalyst used in the low temperature stage is very sensitive to high temperatures, and could not operate effectively in the high temperature stage. Thus, the bulk of the reaction is carried out at high temperature to recover most of the heat. The gas is then removed at low temperature, where the equilibrium is much more favourable, on the very active but unstable copper catalyst.

The gas mixture now contains about 18% carbon dioxide which is removed by scrubbing the gas with a solution of a base, using one of several available methods. One that is favoured is an organic base (in the carbon dioxide absorber), a solution of an ethanolamine, often 2,2'-(methylimino)bis-ethanol (N-methyl diethanolamine). The carbon dioxide is released on heating the solution in the carbon dioxide stripper). Much of it is liquefied and sold, for example, for carbonated drinks, as a coolant for nuclear power stations and for promoting the growth of plants in greenhouses . The last traces of oxides of carbon are removed by passing the gases over a nickel catalyst at 600 K: This process is known as methanation. A gas is obtained of typical composition: 74% hydrogen, 25% nitrogen, 1% methane, together with some argon. Hydrogen from naphtha If naphtha is used as the feedstock, an extra reforming stage is needed.

The naphtha is heated to form a vapour, mixed with steam and passed through tubes, heated at 750 K and packed with a catalyst, nickel supported on a mixture of aluminium and magnesium oxides. The main product is methane together with oxides of carbon, and is then processed by steam reforming, as if it was natural gas, followed by the shift reaction. Hydrogen from coal If coal is used, it is first finely ground and heated in an atmosphere of oxygen and steam. Some of the coal burns very rapidly in oxygen (in less than 0.1 s) causing the temperature in the furnace to rise and the rest of the coal reacts with the steam: The gas emitted contains ca 55% carbon monoxide, 30% hydrogen, 10% carbon dioxide and small amounts of methane and other hydrocarbons. This mixture is treated by the shift reaction. The main problems of using coal includes the large amounts of sulfur dioxide and trioxide generated in burning coal and the significant amounts of other impurities such as arsenic and bromine, all of which are very harmful to the atmosphere and all of which are severe poisons to the catalysts in the process.

There is also a massive problem with disposal of the ash. (b) The manufacture of ammonia (The Haber Process) The heart of the process is the reaction between hydrogen and nitrogen in a fixed bed reactor . The gases, in stoichiometric proportions, are heated and passed under pressure over a catalyst (Figure 3). Figure 3 A diagram illustrating a conventional synthesis reactor (a converter). The proportion of ammonia in the equilibrium mixture increases with increasing pressure and with falling temperature (Le Chatelier's Principle). Quantitative data are given in Table 1. To obtain a reasonable yield and favourable rate, high pressures, moderate temperatures and a catalyst are used. Table 1 Percentage, by volume, of ammonia in the equilibrium mixture for the reaction between nitrogen and hydrogen at a range of temperatures and pressures. A wide range of conditions are used, depending on the construction of the reactor. Temperatures used vary between 600 and 700 K, and pressures between 100 and 200 atmospheres. Much work is being done to improve the effectiveness of the catalyst so that pressures as low as 50 atmospheres can be used.

As the reaction is exothermic, cool reactants (nitrogen and hydrogen) are added to reduce the temperature of the reactors (Figure 3). The ammonia is usually stored on site (step 7) and pumped to another part of the plant where it is converted into a fertilizer (urea or an ammonium salt). However it is sometimes transported by sea (Figure 4) or by road, to be used in another plant. Figure 4 In a plant in Western Australia, the ammonia is transferred by pipeline to a nearby harbour (Figure 2, step 8) and transported by ship. This one is carrying about 40 000 tonnes of liquefied ammonia. By kind permission of Yara International ASA. The original catalyst that Haber used was Fe3O4, which was reduced by the reactant, hydrogen, to iron. Much work was done to improve the catalyst and it was found that a small amount of potassium hydroxide was effective as a promoter . Recently research has been focussed on finding even more effective catalysts to enable the process to take place at lower pressures and temperatures. Ruthenium on a graphite surface is a promising one. Figure 5 The production of ammonia.

Postscript The Haber Process is of such importance to our lives that it has figured in three Nobel Prizes in chemistry, all to German scientists, over a period of nearly 90 years, a remarkable record. The first was given in 1918, to Fritz Haber, the chemist who developed the process in the laboratory. The second was to Carl Bosch, whose brilliant engineering skills made the process viable on a massive scale, but who waited until 1931 for his award. In 2007, Gerhard Ertl was awarded the Prize for his work on catalysis of gaseous reactions on solids . Among the wide range of reactions he studied, he gained evidence for the adsorption of nitrogen molecules and hydrogen molecules on the surface iron and that these adsorbed molecules dissociate into atoms. These atoms then join up in stages to form the ammonia molecule. It must be remembered that the conditions used in these studies (at less than 10-10 atm) are very different from the conditions used in industry, ca 150 atm. Summary The whole process of producing ammonia from methane is summarized in Figure 5. If coal or naphtha is the feedstock, extra processes are needed.

Naphtha is converted into methane and oxides of carbon before going into the primary reformer and thence to the shift reaction. Coal is also converted into hydrogen and carbon oxides and this mixture then undergoes the shift reaction. The Haber Process - Nitrogen Fertilizer from the Air Putting the Carbon Back in your Soil The Haber Process Nitrogen Fertilizer from the Air The Haber Process, also called the Haber-Bosch Process, is a complex chemical procedure that takes nitrogen from the air and under high pressures and temperatures combines it with hydrogen to produce ammonia. This ammonia is the base of the synthetic nitrogen fertilizers increasingly used around the world today. Nitrogen, a key component of all proteins, DNA, and RNA, is vital to life here. Plants can only use fixed nitrogen and the lack of fixed nitrogen is often the limiting factor in an ecosystem or for crops. Our ability to fix nitrogen ourselves seemed to be a perfect solution. However, the enormous blessings of the Haber Process are balanced by some pretty serious curses. On the blessing side synthetic nitrogen fertilizer produced by the Haber Process is credited with feeding a third to half the present world population.

In fact about half the nitrogen in each of our bodies is there thanks to the Haber Process. On the curse side we have several issues including: Serious imbalances to the nitrogen cycle. High fossil fuel energy inputs. Negative effects on soil organisms and soil organic matter. Excess runoff cause ocean dead zones. Major component of weapons including all those roadside bombs. A Very Short History Through the 1800's, heavily populated countries in Europe such as Germany and Britain used guano from islands off the coast of Peru and salt petre from Chile as a source of natural nitrogen fertilizer. This trade was much like our current oil trade in importance both for food security and politically. By the end of the 19th century these sources of fixed nitrogen were running out. Scientists were challenged to come up with a process of making fixed nitrogen from the nitrogen in the air. The winner of the race was Fritz Haber. His tabletop contraption used high pressure, heat and a catalyst to slowly, drop by drop, synthesize a cup of ammonia over the course of two hours. The German chemical company BASF bought the process and Carl Bosch was charged with upscaling it to factory size, a feat he managed by 1913.

The process, often called the Haber Bosch Process requires high pressures, 200 - 300 atmospheres, and high temperatures, 400 to 550 degrees C. For a brief time the first factory turned out fertilizer. With onset of WW1 Germany needed munitions for its war effort. The factory was seconded to munitions and is credited or blamed with greatly prolonging WW1 and making WW2 possible. These synthetic nitrogen fertilizers are explosive. This photo shows the aftermath of the 1921 Oppau explosion that took place at the BASF ammonium nitrate factory there.Between 500 and 600 were killed and 2000 injured. Impacts of the Haber Process on Farming, Gardening and the World More Food - More People It wasn't until the 1950's, after the first and second world wars, that the Haber process really started to affect farming. Ammonia stocks, diverted in wartime to make bombs and bullets, started being used to produce the synthetic nitrogen fertilizers used everywhere today. Increased food production has allowed the human population to swell to its current almost 7 billion size.

Estimates are that a third of the earth's human population is fed thanks to the Haber process. Many people believe that to stop or limit the use of synthetic fertilizers would lead to mass starvation. Balance in the Nitrogen Cycle Our impact on the nitrogen cycle is actually bigger than our effect on the carbon cycle. We convert more N2 gas into fixed reactive forms than all the Earth's processes combined. Synthetic nitrogen fertilizer production, vehicle exhaust emissions and even growing nitrogen fixing crops all fix nitrogen. We add an estimated 140 million tons of fixed nitrogen of which roughly 80-100 million tons is from the Haber process chemical plants. If you study scientific material on the nitrogen cycle you can't help but notice that molecules of nitrogen that get fixed can also get unfixed. Nitrogen fixing bacteria are balanced with denitrifying counterparts in roughly equal proportions. Over time there might be a slow increase in fixed nitrogen but in the natural world the shift is very slow. The Haber process contributions of reactive nitrogen are not offset by denitrification. A group of scientists led by Johan Rockstrom explored the questions of planetary boundaries in an article recently published by Nature. The usual climate change issue was of course reported.

What might surprise you though is that while we have only just slipped past the edge of the CO2 boundary, the nitrogen cycle disruption was deemed to be about three times over the safe limit. This graphic illustrates the extent to which we may be overstepping the nitrogen boundary. These scientists suggest we reduce our fixed nitrogen additions to about 35 million tons or roughly a third of what we are currently slinging onto fields and belching out of cars and factories. Haber Process is an Energy Glutton Given that the Haber process requires temperatures of 400 - 550C and pressures of 200 - 300 atmospheres it's not surprising that it uses a lot of energy. Manufacture of nitrogen fertilizers uses about 5% of the world's natural gas production, equivalent to 1-2% of the world's annual energy consumption. While natural gas is among the more plentiful fossil fuels, this level of use is not sustainable in the long term. As well, as oil reserves worldwide dwindle, prices will rise putting fertilizers some farmers have come to depend on out of reach for the poor.

Negative Effects on Soil and Soil Community A balanced diet for soil organisms in simple terms is about 20 parts carbon for 1 part nitrogen. Adding nitrogen fertilizer to soil usually creates a surplus of nitrogen. To balance their diet the soil microbes go after carbon reserves in the soil - first the easily digested stuff and then the humus. The organic matter in the soil is lost. The soil food web, that intricate web of civilization living in the soil, is starved and begins to die. The functions they performed, whether they be fixing nitrogen, making phosphorus stores available or defending crops against various pests and diseases, are also lost. As the natural system breaks down the soil organisms who survive tend to be the bad guys. Their usual enemies are gone. Erosion of topsoil increases; pests and diseases increase. Not a good scenario. Farmers, to maintain yields, turn to more chemical solutions. More fertilizers, more herbicides and pesticides... and more trouble for the community making it's home in the dirt. As well, recent research at the Center for Ecology and Evolutionary Biology at the University of Oregon by Dr. Jennifer Fox, found that communication links between plants and nitrogen fixing bacteria are damaged by use of excessive synthetic nitrogen fertilizer.

This leads to less nitrogen fixed naturally and increased reliance on synthetics. Fertilizer Runoff and Pollution Plants can absorb the nutrients available through chemical fertilizers. However, often the dose is too much too fast, and much of the fertilizer is wasted. As little as 10% may be actually absorbed by the crop and rarely more than 50% of the fertilizer applied makes it into the crop grown. The nitrogen fertilizer tends to be especially mobile in the soil. It passes right by the plants and off into ground water polluting wells or is eroded into surface streams. The huge ocean dead zones are a result of nitrogen and other fertilizers and farm chemicals ending up in watersheds bound for the oceans. Bombs and Bullets There is still a good chunk of Haber process nitrogen that winds up in war zones. In fact many of those roadside bombs terrorizing people in various parts of the world are made in part from Nitrogen fertilizer synthesized with this process. Hope - Compost and Organic Matter Restores Soils The synthetic nitrogen fertilizers brought to us complements of the Haber process seemed at first a boon to humanity. Synthetic chemical fertilizers have only been commonly available in the last 50 years.

Using synthetic fertilizers has become so routine farmers who use them are considered conventional while those using the techniques thousands of years old are called organic. Given that so little of the Haber process fertilizer actually makes it into crops (and so much into the ocean) it seems to me that we may be able to feed the world better with a lot less chemical additions if we pay a lot more attention to biology. On a personal scale making and using compost is a step that restores soil and their communities. Can that be a bad thing? Want more information? These books will let you delve more deeply into this topic. The Alchemy of Air is Fritz Haber's biography. I'd balance it with one or more of the other books to explore the hope of soil biology. Canadians can access the Amazon.ca store from the second set of links. Check Out These Pages The Haber Process for the manufacture of ammonia THE HABER PROCESS This page describes the Haber Process for the manufacture of ammonia from nitrogen and hydrogen, and then goes on to explain the reasons for the conditions used in the process. It looks at the effect of temperature, pressure and catalyst on the composition of the equilibrium mixture, the rate of the reaction and the economics of the process.

Important: If you aren't sure about using Le Chatelier's Principle or about the effect of changing conditions on rates of reaction you should explore these links before you go on. When you are reading this page, if you find that you aren't understanding the effect of changing one of the conditions on the position of equilibrium or on the rate of the reaction, come back and follow up these links. A brief summary of the Haber Process The Haber Process combines nitrogen from the air with hydrogen derived mainly from natural gas (methane) into ammonia. The reaction is reversible and the production of ammonia is exothermic. A flow scheme for the Haber Process looks like this: Some notes on the conditions The catalyst The catalyst is actually slightly more complicated than pure iron. It has potassium hydroxide added to it as a promoter - a substance that increases its efficiency. The pressure The pressure varies from one manufacturing plant to another, but is always high. You can't go far wrong in an exam quoting 200 atmospheres. Recycling At each pass of the gases through the reactor, only about 15% of the nitrogen and hydrogen converts to ammonia.

(This figure also varies from plant to plant.) By continual recycling of the unreacted nitrogen and hydrogen, the overall conversion is about 98%. Explaining the conditions The proportions of nitrogen and hydrogen The mixture of nitrogen and hydrogen going into the reactor is in the ratio of 1 volume of nitrogen to 3 volumes of hydrogen. Avogadro's Law says that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. That means that the gases are going into the reactor in the ratio of 1 molecule of nitrogen to 3 of hydrogen. That is the proportion demanded by the equation. In some reactions you might choose to use an excess of one of the reactants. You would do this if it is particularly important to use up as much as possible of the other reactant - if, for example, it was much more expensive. That doesn't apply in this case. There is always a down-side to using anything other than the equation proportions. If you have an excess of one reactant there will be molecules passing through the reactor which can't possibly react because there isn't anything for them to react with. This wastes reactor space - particularly space on the surface of the catalyst.

The temperature Equilibrium considerations You need to shift the position of the equilibrium as far as possible to the right in order to produce the maximum possible amount of ammonia in the equilibrium mixture. The forward reaction (the production of ammonia) is exothermic. According to Le Chatelier's Principle, this will be favoured if you lower the temperature. The system will respond by moving the position of equilibrium to counteract this - in other words by producing more heat. In order to get as much ammonia as possible in the equilibrium mixture, you need as low a temperature as possible. However, 400 - 450°C isn't a low temperature! Rate considerations The lower the temperature you use, the slower the reaction becomes. A manufacturer is trying to produce as much ammonia as possible per day. It makes no sense to try to achieve an equilibrium mixture which contains a very high proportion of ammonia if it takes several years for the reaction to reach that equilibrium. You need the gases to reach equilibrium within the very short time that they will be in contact with the catalyst in the reactor.

The compromise 400 - 450°C is a compromise temperature producing a reasonably high proportion of ammonia in the equilibrium mixture (even if it is only 15%), but in a very short time. The pressure Equilibrium considerations Notice that there are 4 molecules on the left-hand side of the equation, but only 2 on the right. According to Le Chatelier's Principle, if you increase the pressure the system will respond by favouring the reaction which produces fewer molecules. That will cause the pressure to fall again. In order to get as much ammonia as possible in the equilibrium mixture, you need as high a pressure as possible. 200 atmospheres is a high pressure, but not amazingly high. Rate considerations Increasing the pressure brings the molecules closer together. In this particular instance, it will increase their chances of hitting and sticking to the surface of the catalyst where they can react. The higher the pressure the better in terms of the rate of a gas reaction. Economic considerations Very high pressures are very expensive to produce on two counts. You have to build extremely strong pipes and containment vessels to withstand the very high pressure. That increases your capital costs when the plant is built. High pressures cost a lot to produce and maintain.

That means that the running costs of your plant are very high. The compromise 200 atmospheres is a compromise pressure chosen on economic grounds. If the pressure used is too high, the cost of generating it exceeds the price you can get for the extra ammonia produced. The catalyst Equilibrium considerations The catalyst has no effect whatsoever on the position of the equilibrium. Adding a catalyst doesn't produce any greater percentage of ammonia in the equilibrium mixture. Its only function is to speed up the reaction. Rate considerations In the absence of a catalyst the reaction is so slow that virtually no reaction happens in any sensible time. The catalyst ensures that the reaction is fast enough for a dynamic equilibrium to be set up within the very short time that the gases are actually in the reactor. Separating the ammonia When the gases leave the reactor they are hot and at a very high pressure. Ammonia is easily liquefied under pressure as long as it isn't too hot, and so the temperature of the mixture is lowered enough for the ammonia to turn to a liquid. The nitrogen and hydrogen remain as gases even under these high pressures, and can be recycled. Questions to test your understanding If this is the first set of questions you have done, please read the introductory page before you start.

You will need to use the BACK BUTTON on your browser to come back here afterwards. BBC - GCSE Bitesize: Ammonia and the Haber process Ammonia and the Haber process Ammonia Ammonia (NH3) is a compound [compound: A compound is a substance formed by the chemical union (involving bond formation) of two or more elements. ] of nitrogen and hydrogen. It is a colourless gas with a choking smell, and a weak alkali which is very soluble in water. Ammonia is used to make fertilisers, explosives, dyes, household cleaners and nylon. It is also the most important raw material in the manufacture of nitric acid. Ammonia is manufactured by combining nitrogen and hydrogen in an important industrial process called the Haber process. An ammonia production plant. Photo courtesy of WMC Resources Ltd The Haber process The raw materials for this process are hydrogen and nitrogen. Hydrogen is obtained by reacting natural gas - methane - with steam, or through the cracking of oil. Nitrogen is obtained by burning hydrogen in air. Air is 80 per cent nitrogen; nearly all the rest is oxygen.

When hydrogen is burned in air, the oxygen combines with the hydrogen, leaving nitrogen behind. Nitrogen and hydrogen will react together under these conditions: a high temperature - about 450ºC a high pressure - about 200 atmospheres (200 times normal pressure) an iron catalyst Haber-Bosch Process - Chemistry Information The Haber process or Haber-Bosch process is the primary industrial method used to make ammonia or fix nitrogen . The Haber process reacts nitrogen and hydrogen gas to form ammonia: N2 + 3 H2 → 2 NH3 (ΔH = −92.4 kJ·mol−1) History of the Haber Process Fritz Haber , a German chemist, and Robert Le Rossignol, a British chemist, demonstrated the first ammonia synthesis process in 1909. They formed ammonia drop by drop from pressurized air. However, the technology did not exist to extend the pressure required in this tabletop apparatus to commercial production. Carl Bosch , an engineer at BASF, resolved the engineering problems associated with industrial ammonia production. BASF's German Oppau plant commenced ammonia production in 1913.

How the Haber-Bosch Process Works Haber's original process made ammonia from air. The industrial Haber-Bosch process mixes nitrogen gas and hydrogen gas in a pressure vessel that contains a special catalyst to speed the reaction. continue reading below our video What are the Seven Wonders of the World From a thermodynamic standpoint, the reaction between nitrogen and hydrogen favors the product at room temperature and pressure, but the reaction does not generate much ammonia. The reaction is exothermic; at increased temperature and atmospheric pressure, the equilibrium quickly switches the other direction. So, the catalyst and increased pressure are the scientific magic behind the process. Bosch's original catalyst was osmium, but BASF quickly settled upon a less-expensive iron-based catalyst, which is still in use today. Some modern processes employ a ruthenium catalyst, which is more active than the iron catalyst. Although Bosch originally electrolyzed water to obtain hydrogen, the modern version of the process uses natural gas to obtain methane, which is processed to get hydrogen gas. It is estimated that 3-5% of the world's natural gas production goes toward the Haber process. The gases pass over the catalyst bed multiple times, since conversion to ammonia is only around 15% each time.

By the end of the process, about 97% conversion of nitrogen and hydrogen to ammonia is achieved. Importance of the Haber Process Some people consider the Haber process to be the most important invention of the past 200 years! The primary reason the Haber process is important is because ammonia is used as a plant fertilizer, enabling farmers to grow enough crops to support an ever-increasing world population. The Haber process supplies 500 million tons (453 billion kilograms) of nitrogen-based fertilizer annually, which is estimated to support food for a third of the people on Earth. There are negative associations with the Haber process, too. In World War I, the ammonia was used to produce nitric acid to manufacture munitions. Some argue the population explosion, for better or worse, would not have happened without the increased food available because of the fertilizer. Also, the release of nitrogen compounds has had a negative environmental impact. References Enriching the Earth: Fritz Haber, Carl Bosch, and the Transformation of World Food Production, Vaclav Smil (2001) ISBN 0-262-19449-X.

US Enivronmental Protection Agency: Human Alteration of the Global Nitrogen Cycle: Causes and Consequences by Peter M. Vitousek, Chair, John Aber, Robert W. Howarth, Gene E. Likens, Pamela A. Matson, David W. Schindler, William H. Schlesinger, and G. David Tilman ChemEngineering - Ammonia production Ammonia production Export (PDF) Ammonia production facilities provide the base anhydrous liquid ammonia used predominantly in fertilizers supplying usable nitrogen for agricultural productivity. Ammonia is one of the most abundantly-produced inorganic chemicals. There are literally dozens of large-scale ammonia production plants throughout the industrial world, some of which produce as much as 2,000 to 3,000 tons per day of anhydrous ammonia in liquid form. The worldwide production in 2006 was 122,000,000 metric tons. [1] China produced 32.0% of the worldwide production followed by India with 8.9%, Russia with 8.2%, and the United States with 6.5%.

Without such massive production, our agriculturally-dependent civilization would face serious challenges. History Before the start of World War I, most ammonia was obtained by the dry distillation of nitrogenous vegetable and animal products; the reduction of nitrous acid and nitrites with hydrogen; and the decomposition of ammonium salt]s by alkaline hydroxides or by quicklime, the salt most generally used being the ammonium chloride (sal-ammoniac). The Haber process, which is the production of ammonia by combining hydrogen and nitrogen, was first patented by Fritz Haber in 1908. In 1910, Carl Bosch, while working for the German chemical company BASF, successfully commercialized the process and secured further patents. It was first used on an industrial scale by the Germans during World War I. Since then, the process has often been referred to as the "Haber-Bosch process". Modern production process A typical modern ammonia-producing plant first converts natural gas (i.e., methane) or LPG Block-flow diagram of ammonia synthesis process (liquified petroleum gases such as propane and butane) or petroleum naphtha into gaseous hydrogen. The method for producing hydrogen from hydrocarbons is referred to as "steam reforming".

[2] [3] [4] [5] The hydrogen is then combined with nitrogen to produce ammonia. [6] [7] Starting with a natural gas feedstock, the processes used in producing the hydrogen are: The first step in the process is to remove sulfur compounds from the feedstock because sulfur deactivates the catalysts used in subsequent steps.

Sulfur removal requires catalytic hydrogenation to convert organic sulfur compounds (RSH) in the feedstocks to gaseous hydrogen sulfide (H2S): H2 + RSH → RH + H2S (gas) The gaseous hydrogen sulfide is then passed through a reactor containing zinc oxide (ZnO) where it is converted to solid zinc sulfide (ZnS) in a batch process (i.e., when allof the zinc oxide has been converted to zinc sulfide, it is removed and replaced with fresh zinc oxide): H2S + ZnO → ZnS + H2O Catalytic steam reforming of the sulfur-free methane (CH4) feedstock is then used to form carbon monoxide (CO) plus hydrogen (H2): CH4 + H2O → CO + 3H2 The next step then uses catalytic shift conversion to convert the carbon monoxide to carbon dioxide (CO2) and more hydrogen: CO + H2O → CO2 + H2 The carbon dioxide is then removed either by absorption in aqueous ethanolamine solutions or by adsorption in pressure swing adsorbers (PSA) using proprietary solid adsorption media.

The final step in producing the hydrogen is to use catalytic methanation to remove any small residual amounts of carbon monoxide or carbon dioxide from the hydrogen by converting them into methane: CO + 3H2 → CH4 + H2O CO2 + 4H2 → CH4 +2H2O To produce the desired end-product ammonia, the hydrogen is then catalytically reacted with nitrogen (N2) derived from process air to form anhydrous liquid ammonia (NH3). This step is known as the "ammonia synthesis loop" (also referred to as the Haber-Bosch process): 3H2 + N2 → 2NH3 The nitrogen required for the Haber-Bosch process is obtained in either of these two ways: The catalytic steam reformer has two sections: a primary reformer and a secondary reformer.

By the introduction of air into the secondary reformer, the air reacts with some of the hydrogen produced in the primary reformer to form nitrogen and water: [8] H2 + air → N2 + H2O By the low-temperature distillation of air The steam reforming, shift conversion, carbon dioxide removal and methanation steps each operate at absolute pressure s of about 25 to 35 bar, and the ammonia synthesis loop operates at absolute pressures ranging from 60 to 180 bar depending upon which proprietary design is used. There are many engineering and construction companies that offer proprietary designs for ammonia synthesis plants. Haldor Topsoe of Denmark, Technip of France, Uhde GmbH of Germany, and Kellogg Brown & Root of the United States are among the most experienced companies in that field. Uses of ammonia About 80% or more of the ammonia produced is used for fertilizing agricultural crops in the form of aqua ammonia (an aqueous solution of ammonia), ammonium sulfate (NH4)2SO4, ammonium phosphate (NH4)3PO4, ammonium nitrate NH4NO3 and urea (NH2)2CO.

Some anhydrous liquid ammonia is also used directly as a fertilizer. Ammonia is also used for: Manufacture of nitric acid (HNO3) Manufacture of nylon and other polyamides Refrigerant in household, commercial and industrial refrigeration systems Manufacture of dyes Cleaning solutions Safety concerns with ammonia There are several safety-related concerns with anhydrous ammonia and with aqueous solutions of ammonia. It is a respiratory irritant that is a highly hazardous chemical. [9] Release could take place through a simple industrial or transportation accident, a deliberate release caused by terrorists, or by improper handling by those using it in the illegal synthesis of methamphetamines. Storage tanks on farms used for dispensing ammonia as fertilizer are referred to as ''nurse'' tanks and contain approximately 2,500 pounds (1134 kg) of anhydrous ammonia, so any farm with four or more nurse tanks needs to assess its safety. In fact, the U.S. Environmental Protection Agency (U.S.

EPA) mandates the performance of an "Offsite Consequence Analysis" (OCA) as part of their "Risk Management Plan" (RMP) requirements for any facility that stores more than 10,000 pounds (4,536 kg) of anhydrous liquid ammonia or 20,000 pounds (9,072 kg) of aqueous solutions of ammonia. [10] The RMP requirements apply for ammonia refrigeration systems or any other ammonia storage facilities as well as farms. The U.S. Occupational Health & Safety Administration (OSHA) has mandated very similar requirements as part of their "Process Safety Management" (PSM) regulations for hazardous chemicals. [11] In transport, ammonia containers must have proper hazardous material placards and, if the pertinent threshold quantity is exceeded, may need additional safeguards such as reporting the shipment to industry monitoring services such as CHEMTREC [12] or additional local agencies. There may be restrictions on transporting hazardous materials through tunnels, or possibly streets in high-density areas. The U.S.

Department of Homeland Security (DHS), citing its major concern as toxic release, lists anhydrous ammonia, or mixtures containing at least 1 percent ammonia, when stored in quantities of 10,000 pounds or more, as a chemical of interest, which falls under the Risk for Chemical Facility Anti-Terrorism Standards (CFATS) regulations and guidance. [13] Organizations that store or transport more than the threshold quantity of 10,000 pounds, or believe they are at a higher than normal risk, should use the Chemical Security Assessment Tool. [14] The U.S. EPA has issued an additional safeguards document, with special emphasis on the theft of ammonia. [15] Image credits The credit for the photo in the lede of this article is: © Courtesy of Burrup Fertilisers Pty Ltd. The schematic diagram in this article was created by Milton Beychok and is released into the Public Domain.

References ^ Gary Maxwell (2004),Synthetic Nitrogen Products: A Practical Guide to the Products and Processes, 1st Edition, Springer Publishing, ISBN 0-306-48225-8 ^ Samuel Strelzoff (1987), Technology and Manufacture of Ammonia, Krieger Publishing, ISBN 0-89464-250-2} ^ Martyn V. Twygg (1989), Catalyst Handbook, 2nd Edition, Oxford University Press, ISBN 1-874545-36-7 ^ Editor: John J. Mcketta (1993), Chemical Processing Handbook, 1st Edition, CRC Press, ISBN 0-8247-8701-3 ^ Editor: J.R. Jennings (1991), Catalytic Ammonia Synthesis, 1st Edition, Springer Publishing, ISBN 0-306-43628-0 ^ Sami Matar and Lewis F.

Hatch (2001), Chemistry of Petrochemical Processes, 2nd Edition, Gulf Publishing, ISBN 0-8841-5315-0 ^ See references 5 and 6 above Ammonia Production with the Haber- Bosch Process Ammonia Production with the Haber- Bosch Process Nearly all modern commercial production of ammonia is based on the Haber-Bosch synthesis process. The process reaction is given by the equation: N2 + 3 H2 < == > 2 NH3 + Heat This reaction happens in a special high pressure reactor in the presence of a special catalyst, usually a porous iron oxide. The reaction is exothermic which means that energy is released. The equation shows that the nitrogen, hydrogen and the ammonia exist in equilirium which is determined by the conditions existing in the reactor. Typically for ammonia synthesis these condions are: Pressure - about 150 atmospheres Temperatures 370 - 500 oC Definition: Under equilibrium conditions the proportion of reactants and the product of a chemcal reaction are balanced and detrmined by the existing physical contions such as pressure, temperature and concentrations.

Factors influencing the rate of ammonia production (rate of reaction rx): Temperature - because the reaction is exothermic, lowering the temperature in the reactor will increase the yield of ammonia. But this also slows down the reaction therefore for the reason of efficiency in commercial production the temperature is kept as high as possible. Pressure - increasing the pressure will increase the yield of ammonia but there is a limit in pressure for safety reasons. Click here to read about equilibrium conditions for ammonia reaction The Process (Haber-Bosch): In large scale commercial ammonia production plants the feed stock which makes up the reactants are water, methane and air. Through a series of chemical processes the feedstock is converted to nitrogen, hydrogen and carbon dioxide. After removal of the waste carbon dioxide, the remaining mixture of nitrogen and hydrogen is called "synthesis gas". The synthesis gas enters the ammonia reactor, where it undergoes the synthesis reaction shown above. The synthesis reactor consists of a number of beds containing the catalyst through which the reactants pass and are converted to ammonia under equilibrium conditions. Because the reaction is exothermic, heat is removed by coolers (heat exchangers). This ensures that the maximum amount of ammonia is converted.

At the end of the reactor a stream of mixed gas (ammonia + nitrogen + hydrogen) is removed and cooled. The cooled ammonia condenses and is separated from the other gas which is then returned to the reactor for reprocessing. This is a continous process - at the one end ammonia is continuously removed and this is balanced by new "synthesis gas" which is continously added to the reactor. Quarkology - 94B Haber Process 9.5 Industrial 9.4.B - The Haber Process The Haber process, also called the Haber-Bosch process, is the industrial process used to produce ammonia from nitrogen gas and hydrogen gas, over an enriched iron or ruthenium catalyst. Despite the fact that 78.1% of the air we breathe is nitrogen, the gas is relatively unavailable because it is so unreactive: nitrogen molecules are held together by strong triple bonds. It was not until the early 20th century that the Haber process was developed to harness the atmospheric abundance of nitrogen to create ammonia, which can then be oxidized to make the nitrates and nitrites essential for the production of nitrate fertilizer and explosives.

Prior to the discovery of the Haber process, ammonia had been difficult to produce on an industrial scale. The Haber process is important today because the fertilizer generated from ammonia is responsible for sustaining one-third of the Earth's population. It is estimated that half of the protein within human beings is made of nitrogen that was originally fixed by this process, the remainder was produced by nitrogen fixing bacteria and archaea. In this topic students: Evaluate the significance of this discovery at that time in world history Identify some of the uses of ammonia Use Le Chatelier's Principle to predict the effect of changing reaction conditions on an equilibrium system Describe and explain the reaction conditions that maximise both rate and yield in the production of ammonia in the Haber Process Describe the process of producing ammonia on an industrial scale Ammonia Ammonia or azane is a weak base and compound of nitrogen and hydrogen with the formula NH3. It is a colourless gas with a characteristic pungent odour. Ammonia contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to fertilisers. Ammonia is also a building block for the synthesis of explosives and many pharmaceuticals and is also used in many commercial cleaning products.

In 2006, worldwide production was estimated at 146.5 million tonnes Although in wide use, ammonia is both caustic and hazardous. Ammonia is often called anhydrous ammonia because of the absence of water in the material. Because ammonia boils at -33.34°C at atmospheric pressure, the liquid must be stored under high pressure or at low temperature. Household ammonia or ammonium hydroxide is a solution of NH3 in water. Why Make Ammonia? Approximately 83% of industrially produced ammonia is used to make fertilisers. When applied to soil, it helps provide increased yields of crops such as corn and wheat. All plants need nitrogen to make amino acids, proteins and DNA, but the nitrogen in the atmosphere is not in a form that they can use. Nitrogen fixation is the name given to any process by which nitrogen (N2) in the atmosphere is converted into ammonia. Atmospheric nitrogen is relatively inert and it does not easily react with other chemicals to form new compounds. Fixation processes free up the nitrogen atoms from their diatomic form to be used in other ways.

Nitrogen fixation is essential for all forms of life because nitrogen is required to biosynthesize basic building blocks of plants, animals and other life forms such as nucleotides for DNA and RNA and amino acids for proteins. Nitrogen fixation through the Haber process is essential for agriculture and the manufacture of fertiliser as the ammonia acts as the precursor to various ammonium salts such as ammonium nitrate and sulfate that provide usable nitrogen to plants and then to animals. Ammonia also acts directly or indirectly as the precursor to most nitrogen-containing compounds. Virtually all synthetic nitrogen compounds are derived from ammonia. An important derivative is nitric acid. This key material is generated via the Ostwald process by oxidation of ammonia with air over a platinum catalyst at 700–850°C NH3 + 2O2 ⇒ HNO3 + H2O Both nitric acid and ammonia are used in the production of explosives such as TNT and ammonium nitrate. Household ammonia is a solution of ammonia in water (aqueous ammonium hydroxide) and is used as a general purpose cleaner for many surfaces. Because ammonia results in a relatively streak-free shine, one of its most common uses is to clean glass, porcelain and stainless steel.

It is also frequently used for cleaning ovens and soaking items to loosen baked-on grime. Significance in World History In 1912, Europe was on the brink of World War I. Since ammonia was used in the manufacture of explosives and nitrogen-based fertilisers, demand was far greater than actual supply. At the time most of the world's supply of nitrates came from a giant salt petre on the coast of Chile. At the time of WWI the allied forces had blocked the supply route and Germany needed a new source. Haber developed the process for the synthesis of ammonia in response to the need for more nitrogen compounds. The manufacture of fertilisers for continued food production and the continued manufacture of explosives were high priorities for Germany on the brink of war. Haber's contribution overcame supply problems for Germany's war effort and is held by many historians to be responsible for prolonging the war. The synthesis of ammonia is a reversible reaction with the forward reaction being exothermic. N2 (g) + 3H2 (g) ⇔ 2NH3 (g) ΔH = -92 kJ mol-1 Starting materials for the synthesis of ammonia.

Despite the fact that 78.1% of the air we breathe is nitrogen, the gas is relatively unavailable because it is so unreactive: nitrogen molecules are held together by strong triple bonds. It was not until the early 20th century that the Haber process was developed to harness the atmospheric abundance of nitrogen to create ammonia, which can then be oxidized to make the nitrates and nitrites essential for the production of nitrate fertiliser and explosives. Ammonia is synthesised from gaseous nitrogen and hydrogen. The nitrogen is extracted from air by distillation and the hydrogen can by produced by a number of methods including the reaction of super heated steam with methane forming carbon monoxide and hydrogen. Ammonia is a gas at room temperature so to synthesise and collect it the system must be closed. Closed systems bring reversible reactions to equilibrium so it is important to consider the yield of ammonia when the system comes to equilibrium. Under standard conditions of temperature and pressure, the reaction is slow and has a low yield (an equilibrium position favouring the reactants). It was Haber who used the principles of chemical kinetics and equilibrium to ensure the reaction conditions were set so as to maximise both the rate and the yield.

The Haber Process - A Delicate Balancing Act An increase in temperature always increase the rate of a chemical reaction. Increasing the temperature for the synthesis of ammonia will increase the rate of the forward and reverse reactions and ensure the system gets to equilibrium quickly. However, because the forward reaction is exothermic, an increase in temperature will cause a shift toward the reactants and decrease the yield. According to Le Chatelier's principle, when a change is made to a system in equilibrium, the system will shift to partially counteract the change. This reaction shifts in reverse because the reverse reaction is endothermic and will therefore counteract the increased temperature by absorbing some of the heat. Therefore, increasing temperature speeds up the reaction that produces ammonia but it reduces the yield of ammonia. If the reaction were carried out at a low temperature to maximise the yield, the molecules would not have sufficient energy to react. Haber was able to overcome the temperature issue by developing a suitable catalyst for the reaction. The catalyst lowered the activation energy for the reaction and helped increase the rate that had been lowered by the need for a lower temperature to increase the yield. The catalyst developed was a mixture of Fe2O3 and Fe3O4.

The reaction is usually run at about 400°C, this temperature being a compromise between a higher temperature which would reduce the yield of ammonia and a lower temperature which would reduce the rate of the reaction. Pressure is another factor that will affect the yield of ammonia in the Haber Process. An increase in pressure will increase the rate of the reaction because there will be more successful collisions between reactant molecules. It will also cause a shift in the equilibrium toward the production of ammonia. This is because when pressure is raised, the equilibrium will shift to partially counteract the high pressure. The forward shift will use four moles of molecules and form two moles of molecules, reducing pressure. In the process more ammonia is formed and its yield is increased. The favoured pressure is around 250 atm. Industrial Production In 1909, the Haber process was purchased by the German chemical company BASF, which assigned Carl Bosch the task of scaling up Haber's tabletop machine to industrial-level production. The Haber process on an industrial scale showing each of the reaction conditions being maintained to ensure the best rate and yeild of ammonia. The diagram above shows how ammonia is produced via a continuous industrial process designed to ensure that the system never gets to equilibrium.

The reaction vessel must be continually monitored to ensure that optimal conditions are met for the production of ammonia and to ensure a safe environment for workers with minimal waste. The factors that need to be monitored include: Temperature - This needs to be maintained at around 400°C. If it is too high the yield of ammonia drops. If it is too low the reaction rate drops. Pressure - This needs to be maintained at around 250 atm. If it is too high the vessel may not be strong enough to hold the gases and an explosion may occur. If it is too low the yield of ammonia will drop. Reactants - The ratio of nitrogen to hydrogen needs to be maintained at 1:3 otherwise one reactant will be in excess and wasted. These gases are also pressurised to concentrate them before the enter the system. Product Removal - The ammonia gas is liquefied and removed which reduces its concentration and drives the equilibrium to the right increasing the yield. Page Links

Question: Which gas is produced via the 'Haber Process'?
Answer: AMMONIA